Chlorine gas and sodium hydroxide are jointly manufactured almost exclusively by the electrochemical reduction of salt brine solutions. This electrolytic reaction produces these chemicals as co-products, and also produces hydrogen gas. The three principal commercial manufacturing processes for chlorine and sodium hydroxide are the membrane cell, the diaphragm cell and the mercury cell methods. The mercury cell process accounts for approximately 15% of the total production of chlorine and sodium hydroxide. See Chemical and Process Technology Encyclopedia, by McGraw-Hill, edited by Douglas M. Considine, 1974, at page 275.
All three manufacturing processes require substantial amounts of electrical energy to accomplish the reduction. In fact, depending upon local electric power availability, generally over 50% of the cost to manufacture these important industrial chemicals is attributable to energy requirements alone. Obviously, any reduction in such energy needs has a material impact on total manufacturing costs and would be an important development in this basic industry. The improvement in the process described herein represents an approximate 2% to 10% reduction in energy requirements over the older and prevailing process for a given volume of product. The significance of such a development is reflected in the amount of chlorine produced annually in this country, reportedly in excess of 10 million pounds per year (Id., at page 273).
In either the mercury cell or diaphragm cell methods, an anode and cathode are immersed in a brine solution. When a current is directed through the electrolyte, chlorine is produced at the anode (typically made of graphite) while sodium hydroxide is produced near the cathode (Id., at page 275).
Production of chlorine by the mercury cell process involves both a brine cycle and a mercury cycle, two continuous and interactive fluid cycles. The process dictates that most modern mercury cells be similar in construction and appearance with an electrolyzer section and a decomposer section to accommodate the two fluid cycles, for instance, see FIG. C-18, Id., at page 276. In the long rectangular electrolyzer section, chlorine gas forms on the anode, bubbles to the top, and is removed. Metallic sodium formed at the mercury cathode immediately amalgamates with the mercury flowing along the bottom under a gravity head. The depleted brine and mercury/sodium amalgam flow concurrently into the next equipment section, the decomposer, where deionized water is added to form sodium hydroxide and hydrogen gas, thus regenerating the mercury for recycle. The concentrated caustic solution is removed as a product. Modern cells may operate in the foregoing manner using more than 300,000 amps and a current density of over 10,000 amps per square meter at the cathode.
Important and notable differences between modern mercury cell designs are in the configuration of the decomposer section where sodium hydroxide is generated and in the technique for adjusting the anode to cathode gap in the electrolyzer section to minimize voltage drop. The voltage across the cell is directly proportional to power requirements (P=VI) and hence, energy consumption (W=VIt) [assuming a constant amperage, I, which enables production of a constant flow rate of product generated over time, t]. The anode section is typically an array of several anodes suspended from an overhead bus running across the electrolyzer above the mercury cathode. For instance, see FIG. C-19, Id., at page 277. To minimize voltage drop, the group of anodes is lowered until there is an incipient short circuit. The anodes are then raised so that the clearance under dynamic conditions is the least practical distance to obtain minimum voltage drop across the cell while allowing for reasonable anode life. Typically, the voltage will be 4.3 volts at about 10,000 amps per square meter in an ordinary modern production cell.
Side reactions at both the cathode and anode can interfere with the primary electrochemical reduction. For example, hypochlorite is a natural by-product of the reaction between chlorine (formed at the anode) and water. This reactant can erode graphite anodes and reduce the yield for a given current. See Operating Manual for A Plant for the Electrolysis of Alkaline Chloride, Friedrich Uhde, GmBH., at pages 21-22. Contaminants in the brine can also cause inefficiencies at the cathode. Magnesium in the feed brine can lead to magnesium hydroxide precipitation at the mercury cathode. The pH in the boundary layer near the cathode is believed to be sufficiently alkaline to cause this precipitation which tends to block the transport of sodium ions, which in turn raises the electrode resistance potential and energy consumption. See Salt Resources and Brine Treatment in Chloralkali Production, by J. H. G. van der Stegen, Akzo Report. 90.00.007, at pages 58-59.
In the invention, it has now been determined that by adding to the cell feed brine a small amount of water soluble polymer containing polar functional groups, the cell voltage can thereby be reduced, and consequently, overall power requirements are reduced for the same volume of chemical production. This small amount of polar polymeric compound in the feed stream is believed to inhibit the formation of interfering precipitates such as magnesium hydroxide and possibly the hydrolysis derivatives of chlorine, such as hypochlorite.
One example of such a polymeric material is polyacrylic acid in an emulsifying solution. This material, once dissolved in an aqueous medium, can be metered into the feed brine in minute amounts and effectuate an approximate 2% to 10% reduction in electrical power consumption for a given product yield as compared with the same cell conditions without the polymer present.